The pH-value is an integral parameter of every (aqueous) solution. It describes to which degree the solution is alkaline or acidic. Over a wide range it is well approximated by: pH=−log10 [H+], wherein [H+] denotes the hydrogen ion concentration of the solution in mol/L. Measuring a pH-value of an aqueous solution is a routine task in the industry and also in laboratories for process control and analysis. However, it could also become interesting for a wider range of applications if the pH-measurement units (sensor plus electronics) become sufficiently inexpensive. For example, there is a large potential for pH-measurement to monitor the quality of (liquid) perishables in the supply chain or even at the customer's himself. Experimental techniques for measuring ion concentrations (in particular pH) can be divided into two classes, non-electrochemical methods, e.g. optical (indicator dyes), catalytic, and swelling of polymers (gels), and electrochemical methods. The latter are widely used for many applications in industry and laboratories. Electrochemical ion concentration sensors rely on the potentiometric principle, i.e. they measure the electrical potential φ at a solid/liquid interface or across a membrane which is a function of the ion concentration to be determined. φ can be calculated from the Nernst equation: φ=kT/(nq) ln(a1/a2), wherein k is the Boltzmann constant, T the absolute temperature in Kelvin, q the elementary charge, n the ionic charge (e.g., n=1 for H3O+, Na+; n=2 for Ca2+), and a1, a2 the respective activities at both sides of the membrane/interface.
Ion concentrations at both sides of the membrane/interface are represented in terms of activities ai=fi*ci with fi being the respective activity coefficient (fi=1 for diluted electrolytes) and ci the respective ion concentration. According to the Nernst equation the electrode potential is a logarithmic function of the ion activity on one side of the membrane/interface if the activity on the other side is kept constant. Depending on the type of ion described by “a”, the sensor is sensitive to H3O+-ions, Na+-ions, Ca2+-ions, etc.
All major pH (ion) measurement electrodes operate according to the principle described above, including the well-known glass electrodes (different glass compositions have been developed that are sensitive to pH, pNa, pK, etc., respectively), antimony electrodes, ISFET's (Ion Sensitive Field Effect Transistors) and EIS capacitors (Electrolyte Insulator Semiconductor capacitors; here the flat-band voltage is a function of the pH/pNa/pK/etc of the electrolyte).
In order to measure the potential difference a reference electrode is needed; for the ISFETS and EIS devices the reference electrode defines the electrolyte potential to set the operating point and do the measurement. The potential of the reference electrode with respect to the electrolyte potential must remain constant irrespective of the electrolyte composition. Besides the standard hydrogen electrode the Ag/AgCl electrode is the most well-known reference electrode. It consists of a chlorinated silver wire in contact with a well defined electrolyte (often 3 mol/L KCl). Galvanic contact between the analyte and the electrolyte is established via a diaphragm, such as a porous frit from glass or ceramics. During operation the electrolyte must continuously flow out of the reference electrode into the analyte. Other reference electrodes, e.g. calomel (based on mercury) or Tl/TlCl electrodes, are used for specific applications, e.g. at elevated temperatures. Their principle is the same as for the Ag/AgCl electrode, in particular with respect to the use of liquid electrolyte and contact via a diaphragm.
The problem with the known electrochemical sensors is that they require a reference electrode in order to determine the charged particle concentration from a measured potential (difference). Using reference electrodes, and in particular accurate reference electrodes, involves all kinds of difficulties such as the following:                Electrolyte outflow in a reference electrode through the diaphragm is essential. That means the electrolyte needs to be refilled regularly. Moreover, the pressure conditions must be such that the outflow is guaranteed, i.e. the pressure in the analyte cannot be higher than in the reference electrode (otherwise the analyte enters the reference electrode and changes its potential, which is called reference electrode poisoning;        Clogging of the diaphragm of the reference electrode causes measurement errors (depending on the application regular cleaning is needed);        Most reference electrodes have rather large dimensions, which makes it difficult/impossible to integrate them into a miniaturized device. Some miniature reference electrodes exist but they have a limited lifetime (because reference electrolyte cannot be refilled);        Reference electrodes have a limited temperature range, e.g., for high temperatures a Tl/TlCl electrode must be used; and        Some reference electrodes may react to other environmental parameters, for example, the silver in Ag/AgCl electrodes is light sensitive.        
Even pseudo-reference electrodes suffer from several disadvantages, such as:                complex (expensive) integration, corrosion, interface leakage, food and bio-compatibility issues.        